Chemistry for Life: Exam 1

Elements of Life

The Atom Top number – Mass number – total number of protons and neutrons Bottom number – Atomic (proton) number – number of protons Number of neutrons = top no. – bottom no. = mass number – atomic number No. of electrons = no. of protons, for neutral atoms (no overall charge) Negative ions have more electrons than protons Positive ions have less electrons than protons No. of protons – no. of electrons = charge on ion Isotopes of an element are atoms with the same number of protons but different numbers

of neutrons Chemical properties decided by number and arrangement of electrons Isotopes have same configuration of electrons so same chemical properties Physical properties depend on mass, and as the mass changes for isotopes, their

physical properties change

Atomic Models Ancient Greeks – indivisible particles Start of 19th century – John Dalton – solid spheres, different spheres made up different

elements 1897 – J J Thompson discovered electrons from charge/mass measurements – model

changed from ‘solid sphere’ to ‘plum pudding’ (Figure 1) 1909 – Ernest Rutherford, Hans Geiger, Ernest Marsden – Gold foil experiment where

alpha particles (positive) are fired at thin sheet of gold. Expected most alpha particles to be deflected slightly by positive ‘pudding’, but most passed through and few were deflected backwards. ‘Nuclear model’ created – tiny, positive nucleus at centre (most of mass) which is surrounded by cloud of negative electrons, most of atom empty space (Figure 2)

Henry Moseley – charge of nucleus increases from one element to next by 1 Rutherford did further experiments and discovered protons which explains this (atoms of

different elements have a different number of protons) James Chadwick – neutrons – nuclei of atoms heavier than they would be with just

protons so something to be discovered with mass but no charge Electrons would spiral down into nucleus causing atom to collapse if electrons existed as

a cloud Neils Bohr – Bohr Model (Figure 3) – electrons exist in fixed orbits/shells with fixed

energies, EM radiation emitted or absorbed when electrons move between shells, radiation has fixed frequency. Shells of an atom hold fixed numbers of electrons, reactivity is due to electrons, full shells means stable/inert/unreactive

Relative Mass Relative mass – mass of an atom compared to carbon-12 Relative atomic mass (Ar) – average mass of an atom of an element on a scale where an

atom of carbon-12 is 12 Relative isotopic mass – mass of an atom of an isotope on a scale where an atom of

carbon-12 is 12 Relative molecular/formula mass (Mr) – average mass of a molecule of formula unit on a

scale where an atom of carbon-12 is 12 Mass spectrometer (Figure 4)

1. Vaporisation – sample injected and is vaporised by electrical heater2. Ionisation – gas particles bombarded by high-energy electrons which knock off

electrons, ionising them and leaving positive ions

3. Acceleration – positive ions accelerated by electric field4. Detection – Mass spectrum produced by measuring time taken for each particle

to reach detector – depends on mass and charge – light highly charged ones are faster

Molecular ion formed when bombarding electrons remove 1 electron, giving highest mass, furthest on right

Bombardment causes some molecules to break up into fragments, which are seen on graph as fragmentation pattern

Nuclear Radiation If unstable, atoms break down to become stable. Instability caused by having too many or

not enough neutrons, or too much energy in nucleusAlpha particles Beta particles Gamma rays

What they are Helium nuclei Fast moving electrons Very short wave EM waves

Stopped by Paper/skin Thin aluminium sheets

Very thick lead

Ionising ability Strong Moderate WeakDeflection in electric field

Slight Large None

When alpha particles hit an atom they transfer some energy to it, they quickly ionise lots of atoms and lose all energy – low penetrating power

Nuclear fusion – 2 small light nuclei combine under high temperature and pressure to make one larger nucleus, happens naturally in stars, huge amounts of energy released

When hydrogen runs out, core’s pressure and temperature rise. In a big enough star it’ll get hot enough to fuse heavier elements.

The Mole and Equations Avagadro’s Constant = 6.02 x 10^23 (L)

Empirical and Molecular Formulae Empirical formula – smallest whole number ratio of atoms in a compound Molecular formula – actual numbers of atoms in a compound

Electron Shells and Atomic SpectraElectron Shell 1 2 3 4Max. no. electrons 2 8 18 32

For a particular element, the frequencies in an emission spectrum are the same as those missing in the absorption spectrum. Each element has different electron arrangement so frequencies of radiation are different so spectrum for each element is unique (figure 5).

Spectra for hydrogen (Figure 6) – when electrons drop down to n=1 line produced in UV part of EM spectrum, n=2 visible light, n=3 infrared

Ionic bonding (between metal and non-metal) Compounds are atoms of different elements bonded together Ionic bonding is when oppositely charged ions become stuck together by electrostatic

attraction – very strong Sodium chloride (NaCl) – giant ionic lattice structure – cubed shaped Conduct electricity when molten or dissolved but not when solid, as ions are fixed in

position by strong ionic bonds High melting points Often dissolve in water (polar)

Covalent bonding (between non-metals) Molecules are groups of atoms bonded together

Both positive nuclei are attracted electrostatically to the shared electrons Figure 7 and 8 – Covalent bonding in molecules and dative covalent bonding Fairly low melting and boiling points – no giant structure to be broken down – only have

to overcome weak intermolecular forces not strong covalent bonds between atoms – simple molecular

Don’t conduct electricity – no free charge carriers Usually insoluble in water van der Waals = intermolecular forces

Bonding Examples M/B point

State at STP Solid conduct elec?

Liquid conduct elec?

Soluble in water?

Ionic NaCl, MgCl2

High Solid No Yes Yes

Simple molecular (covalent)

CO2, I2, H2O

Fairly low

Usually gas, liquid, sometimes solids

No No No (depends on polarity)

Giant Covalent Structures Tetrahedral arrangement Very high melting points Very hard Good thermal conductors – vibrations travel easily through stiff lattices Won’t dissolve Can’t conduct electricity

Metallic Structures Melting points high, affected by: no. of delocalised electrons per atom (the more there

are, the stronger the bonding, the higher the melting point), size of metal ion, and lattice structure

Can be shaped, ductile Good thermal conductors – delocalised electrons pass kinetic energy to each other Good electrical conductors – delocalised electrons carry a current Metals insoluble except liquid metals, due to strength of metallic bonds

Shapes of Molecules Lone pairs repel more than bonding pairs

The Periodic Table Now arranged by increasing proton number 1863 – John Newlands – arranged elements in order of mass, every 8th element similar –

law of octaves – arranged in rows of 7, similar elements lined up in columns. Pattern broke down on 3rd row

1869 – Dmitri Mendeleev – left some gaps where next element didn’t seem to fit Assumptions at the time: all were elements, and all had been discovered 1890s – noble gases discovered Group 1 – more reactive as you go down group Group 7 – less reactive as you go down group Periods 2 and 3 – similar trends in melting and boiling points due to structure and bond

strength More atoms in a molecule mean stronger intermolecular forces Metallic > Giant covalent > Simple molecular > noble gases

Group 2 Elements

Metal + water = metal hydroxide + hydrogen Reactivity increases as you go down group as outer electrons are further from nucleus Oxides and hydroxides are bases Metal oxide + water = metal hydroxides (which dissolve) OH- ions make solutions strongly alkaline Magnesium is exception – slow reaction and hydroxide not very soluble Increasingly alkaline as you go down group due to increase in solubility of hydroxides acid + base = salt + water Carbonates and sulphates decrease in solubility down group carbonate = oxide + carbon dioxide Thermal stability increases down group. Carbonate ions are large anions and are made

unstable by presence of cation. Cation attracts electrons towards, polarising and distorting anion. Greater distortion means less stable. Large cations cause less distortion (lower charge density). So, down the group, more shells of electrons, bigger cations, less distortion caused, so more stable the carbonate anion.

Developing Fuels

Gas Volumes and Entropy Number of moles = volume in dm³ / 24 One mole of any gas occupies same volume (24dm³) at rtp Entropy is a measure of the number of ways the particles can be arranged (measure of

disorder/randomness) Substances like disorder – particles naturally move to give maximum possible entropy.

Gases diffuse to fill all available space – more ways to be arranged in bigger space. When something dissolves, solute particles spread out in solvent, increasing entropy

More particles, higher entropy. A mixture of 2 types of particle has more entropy than same amount of one type

Enthalpy Changes Enthalpy change is the heat energy transferred in a reaction at constant pressure Exothermic reactions give out energy (ΔH-) – temperature of surroundings goes up e.g.

combustion, oxidation of carbohydrates in respiration Endothermic reactions take in energy (ΔH+) – temperature of surroundings goes down

e.g. thermal decomposition, photosynthesis enthalpy change of reaction = energy absorbed to break bonds – energy released in

making bonds Bond breaking – endothermic Bond making – exothermic Bond enthalpies differ because they are averages for a much bigger range of molecules Bond length – distance between 2 nuclei in covalent molecules where attractive and

repulsive forces balance Stronger attraction, higher bond enthalpy, shorter bond length

Hess’s Law Standard enthalpy change of reaction is enthalpy change when reaction occurs in molar

quantities shown in chemical equation under standard conditions in standard states Standard enthalpy change of formation is the enthalpy change when 1 mole of a

compound is formed from its elements in their standard states under standard conditions Standard enthalpy change of combustion is the enthalpy change when 1 mole of a

substance is completely burned in oxygen under standard conditions Hess’s Law – total enthalpy change of a reaction is independent of the route taken,

providing starting and finishing points are the same ΔHr = ΣΔHf (products) – ΣΔHf (reactants)

Measuring Enthalpy Changes

Enthalpy change in a neutralisation reaction:o add a known volume of acid to insulated container and measure temperatureo Add known volume of alkali and record temperature rise (stir for even heating)

Systematic errors are repeated every time you carry out an experiment due to experimental set-up or limitations of equipment:

o some heat absorbed by containero some heat lost to surroundingso some incomplete combustiono some flammable liquid may escape by evaporation

Random errors – no pattern, always happens, difficult to control. Best to repeat experiment and take average of all readings

Reliability – how reproducible results are Accuracy – how close to true value results are

Catalysts A catalyst speeds up the rate of a chemical reaction, but can be recovered chemically

unchanged at the end of the reaction Catalysis means speeding up a chemical reaction by using a catalyst Heterogeneous catalyst examples: iron in Haber process and platinum in catalytic

convertors1. Reactant molecules arrive at heterogeneous catalyst surface and bond with it

(adsorption)2. Bonds between reactant’s atoms are weakened and break-up, forming radicals. Radicals

then get together and form new molecules3. New molecules detach from catalyst (desorption) Adsorption needs to be strong enough to break molecule bonds but not too strong or it

won’t let go of atoms CO poisons solid iron catalyst in Haber process Lead poisons catalytic convertors Heterogeneous catalysts often get poisoned because poison clings to catalyst’s surface

more strongly than reactant does. Catalyst is prevented from getting involved in reaction it should be speeding up

Organic Groups Alkanes – saturated – impossible for carbon to make more than 4 bonds Cycloalkanes have 2 fewer hydrogens than straight-chain alkanes Alkenes – unsaturated – can make more bonds with extra atoms in addition reactions Cycloalkenes have 2 fewer hydrogens than open-chain alkenes Aromatic (arenes) – benzene ring Aliphatic – no ring e.g. alkanes, alkenes Benzene – 6 carbons, 6 hydrogens, 3 double bonds – more stable, less reactive as

double bond electrons are delocalised around carbon ring Ethers have oxygen atom interrupting carbon chain

Isomerism General formula – describes any member of a family of compounds e.g. CnH2n+1OH Molecular formula – actual number of atoms of each element in a molecule, indicating

any functional groups Structural formula – shows atoms carbon by carbon, with attached hydrogens and

functional groups Displayed formula – shows how atoms are arranged and bonds between them Skeletal formula – shows bonds of carbon skeleton only with any functional groups. H

and C atoms not shown Structural isomers – same molecular formula, different structural arrangement of atoms

o different carbon skeleton – different physical properties due to change in shape

o functional group in different place – different physical and sometimes chemical properties

o different functional groups – very different physical and chemical properties

Shapes of Organic Molecules When a carbon atom makes 4 single bonds the molecule is a tetrahedral 3D shape with

bond angles of 109.5º Atoms round a C=C form an equilateral triangle which is trigonal planar (flat) with bond

angles of 120º

Catalysts and Petroleum Crude oil vaporised at 350ºC It then goes into the fractionating column and rises up through the trays. The largest

hydrocarbons don’t vaporise – boiling points are too high so they run to bottom and form residue

As the crude oil vapour rises up through fractionating column it gets cooler. Each fraction condenses at different temperatures due to different chain lengths – fractions drawn off at different levels

Hydrocarbons with lowest boiling points don’t condense and are drawn off as gases at top (Fractional distillation – Figure 9)

Thermal cracking – high temperatures (1000ºC) and high pressure (70 atm), produces lot of alkenes which are used to make polymers e.g. poly(ethene) from ethene

Catalytic cracking – makes motor fuels and aromatic hydrocarbons, uses zeolite catalyst, slight pressure, high temperature (450ºC). Catalyst cuts costs as reaction can be done at lower temperature and pressure and speeds it up

Isomerisation – makes branched-chain isomers. Heating straight-chain alkane with platinum catalyst stuck on inert aluminium oxide. Molecule broken up and put back together as branched-chain isomer. Zeolite then used as molecular sieve to separate isomers. Straight chain molecules go through and are recycled.

Reforming – converts alkanes into Cycloalkanes into arenes – needs catalyst (platinum stuck on aluminium oxide)

Fuel-air mixture squashed by piston and ignited by spark, creating explosion. This drives piston up again, turning crankshaft. 4 pistons work one after other, so engine runs smoothly

Fuels CO molecules bind to same sites on haemoglobin molecules as oxygen The Earth radiates IR radiation out into space, but greenhouse gases absorb some of it,

warming earth Unburnt hydrocarbons and nitrogen oxides react in presence of sunlight to form ground-

level ozone (a component of photochemical smog) which irritates eyes, aggravates respiratory problems

Particulates – tiny particles suspended in the air Sulfur dioxide dissolves in moisture and converted into sulphuric acid Acid rain destroys trees and vegetation, corrodes buildings and statues, kills fish in lakes Sulfur dioxide removed from power station flue gases using calcium oxide Oxygenates added to fuels to reduce CO emissions – fuel combusts fully Governments can change laws and tax pollution more highly Compulsory catalytic converters, MOT emissions test Changing people’s behaviour

Fuels of the Future Short term – reduce pollution, reduce need on fossil fuels Renewable sources e.g. wind, solar, wave power – won’t run out, carbon neutral (except

during manufacture/installation from using machines/vehicles). Not reliable and a lot needed however

Biodiesel/bioethanol – renewable, carbon neutral (except refining and transporting fuel, making fertilisers and powering machinery), expensive, land used for fuel not for food

Nuclear power – no pollution (except in refining and mining uranium ore and decommissioning nuclear plants), radioactive waste, possibility of nuclear disaster

Hydrogen gas burned in modified engine or used in fuel cell (converts hydrogen and oxygen into water to produce electricity)

Takes energy to extract hydrogen from seawater – energy carrier Method of extraction determines how environmentally friendly Hydrogen - difficult to transport/store, flammable, has to be liquefied due to low energy to

volume ratio, need new infrastructure Increasing competition means higher prices, disruption of supply due to political issues Energy efficient, use fossil fuel reserves, financial incentives, renewable energy, nuclear

power

Chemistry of Natural Resources

Elements from the Sea

More Calculations No. of moles = concentration x volume (in cm³) / 1000 Percentage yield never 100% due to not all starting chemicals reacting

fully, or loss of chemicals on filter paper or during transfers between containers

Percentage yield = actual yield / theoretical yield x 100%

Titrations Pipette – measure only one volume of solution Burette – can measure different volumes and allow you to add solution

drop by drop Titrations allow you to find out exactly how much acid is needed to

neutralise a quantity of alkali Measure out alkali using pipette and put in flask, add acid using burette,

swirling flask. Do a rough titration first to find end point, the accurate one where you add acid drop by drop from 2cm³ from endpoint

Universal indicator no good – colour change too gradual These indicators change colour quickly over small pH range

o methyl orange – turns yellow to red when adding acid to alkalio phenolphthalein – turns red to colourless when adding acid to alkali

Electronic Structure The further a shell is from nucleus, the higher its energy and the larger its

principal quantum number Not all the electrons in a shell have exactly the same energy so shells are

divided up into sub-shells that have slightly different energies Sub-shells consist of orbitals which can each hold 2 electrons Chromium and copper donate one of their 4s electrons to the 3d sub-shell

because they’re happier with a more stable full or half-full d sub-shell Groups 4 to 7 can share electrons to form covalent bonds

d block elements (transition metals) tend to lose s and d electrons to form positive ions

Oxidation and Reduction An oxidising agent accepts electrons and gets reduced A reducing agent donates electrons and gets oxidised Combined oxygen oxidation states

o -1 in peroxideso +2 in fluorides OF2

o +1 in O2F2 and H2O2

In metal hydrides hydrogen is -1, otherwise +1 Many metals reduce dilute acids

Electronegativity Electronegativity – the ability to attract the bonding electrons in a covalent

bond Dipole – a difference in charge between 2 atoms caused by a shift in

electron density Molecules will be polar if all polar bonds point in roughly same direction Non-polar molecule if polar bonds point in opposite directions and cancel

each other out, or if lone pairs of electrons on the central atom cancel out dipole created by bonding pairs

Intermolecular forces Much weaker than covalent, ionic or metallic bonds Even though the electron cloud keeps moving and the dipoles keep

changing in instantaneous dipole-induced dipole forces, the overall effect is for the atoms to be attracted to each other

The longer the carbon (alkane) chain, the stronger the van der Waals forces due to more molecular surface area and more electrons to interact

Branched-chain alkanes can’t pack closely together and have smaller molecular surface areas so van der Waals forces reduced

Ionisation Enthalpies The first ionisation enthalpy is the energy needed to remove 1 electron

from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous 1+ ions

Affected by atomic radius (the further the outer electrons are from nucleus the less strongly attracted so IE lower); nuclear charge (the more protons the more strongly attracted the outer electrons are so IE higher); electron shielding (the more inner shells the more shielding from attractive force of nucleus so IE lower)

The 2nd ionisation enthalpy is the energy needed to remove 1 electron from each ion in 1 mole of gaseous 1+ ions to form 1 mole of gaseous 2+ ions

Within each shell, successive IEs increase due to less repulsion amongst remaining electrons and electrons are being removed from an increasingly positive ion so attraction between nucleus and remaining electrons increases

Removing electrons from a paired orbital is easier due to repulsion between the 2 electrons so easier to remove

Group 7 – The Halogens Physical state at rtp:

o fluorine – pale yellow gaso chlorine – green gaso bromine – red-brown liquido iodine – grey solid

Volatility decreases down group Boiling points increase down group due to increasing strength of van der

Waals forces due to increasing size and relative mass of atoms Electronegativity decreases down group as larger atoms attract electrons

less than smaller ones More soluble in organic solvents than water as covalent and non-polar

colour in water colour in hexanechlorine virtually colourless virtually colourlessbromine yellow/orange orange/rediodine brown pink/violet

Less reactive down group as it is harder for larger atoms to attract an electron to form an ion, due to increase electron shielding

Halogens displace and oxidise less reactive halide ions from solution (if halide is below halogen in periodic table)

More about the Halogens Silver ions react with halide ions to form a precipitate – first add dilute

nitric acid to remove ions which might interfere with reaction, then add silver nitrate solution

Fluoride – no precipitate as AgF soluble Chloride – White Bromide – cream Iodide – yellow Electrolysis of halide solutions to extract halogens:

o halogen element produced at anode (+ electrode)o at cathode, hydrogen ions form hydrogen gas

You can only extract chlorine from concentrated sodium chloride solution, if dilute chloride ions aren’t discharged and hang onto their electrons, instead OH- lose their extra electron to form oxygen and water

fluorine can’t be produced by electrolysis of fluoride solutions, even when concentrated OH- discharged instead

Halogen used to make properties used forFluorine PTFE inert, low-friction, non-stick coating

thermally stable on pansHCFCs inert, gas at rtp refrigerantsodium fluoride strengthens tooth

enameltoothpaste

Chlorine PVC electrical insulator electrical wiresbleach kills bacteria water treatment

Bromine medicines, agricultural chemical, flame retardantsIodine Medicines, nutrient

The Chemical Industry Atom economy tells you what proportion of the starting materials end up in

useful products % atom economy = mass of desired product / total mass of reactants x

100

Halogenoalkanes To make 2-chloro-2-methylpropane (a chloroalkane), shake 2-

methylpropan-2-ol (a tertiary alcohol) with concentrated hydrochloric acid at room temperature in separating funnel for 20 mins, releasing pressure due to volatility of product

Allow to settle into layers, run off aqueous lower layer (containing most of unreacted 2-methylpropan-2-ol), leaving impure halogenoalkane

Neutralise excess acid by adding sodium hydrogencarbonate and shaking until no more gas produced, releasing pressure frequently

Run lower layer off, add distilled water and shake, and run off lower layer again (to get rid of remaining inorganic impurities)

Remove remaining water by adding anhydrous sodium sulphate (drying agent) and shake mixture

remove remaining organic impurities (unreacted alcohol) by distilling mixture, collect fraction that boils between 48 and 53C – chloroalkane

More about Halogenoalkanes A nucleophile is an electron-pair donor Iodoalkanes are most reactive Halogenoalkanes due to their lowest bond

enthalpy (weakest carbon-halogen bond strength) 2 experiments to show reactivity:

o put the 3 different haloalkanes in 3 different test tubes, and add silver nitrate (contains water) and ethanol (solvent). Precipitate forms fastest with iodoalkane – most reactive

o warm NaOH (aq) with haloalkanes, add dilute nitric acid to neutralise spare hydroxide ions before adding silver nitrate, or silver nitrate would react with hydroxide ions to form silver oxide precipitate

The Atmosphere

Giant Structures Diamond and silicon(IV) oxide have giant molecular structures

(macromolecular structures) – huge network of covalently bonded atoms Due to carbon and silicon forming 4 strong, covalent bonds In diamond, each C covalently bonded to 4 others, arranged tetrahedrally

in crystal lattice structureo hardo vibrations travel easily through stiff lattice, good thermal conductoro very high melting pointo can’t conduct electricityo won’t dissolve in any solvent

Silicon oxide also arranged tetrahedrally in giant latticeo each silicon atom covalently bonds with 4 O tetrahedrally to form

big crystal latticeo O can only bond with 2 Silicon atomso hard crystalline solid with high melting pointo insoluble in any solvento doesn’t conduct electricity

carbon dioxide – small molecules, each C forming double bond with 2 Os silicon dioxide – each Si bonds singly with 4 O (allowing each O to bond to

another Si) carbon dioxide will dissolve in water, SiO2 won’t

Reaction Rates Particles need to collide in right direction and with certain minimum

amount of kinetic energy (activation enthalpy) in order for a reaction to take place

Maxwell-Boltzmann distribution – graph of no. of molecules of gas with different kinetic energies (as not all molecules in a gas have same amount of energy)

Increase temperature – particles have more energy, collide more often; more molecules have the activation enthalpy

More on Reaction Rates Increasing concentration or pressure – particles closer together on

average, collide more often, more chances to react Increasing surface area – more particles can come in contact with other

reactants Catalysts lower activation enthalpy by providing different way for bonds to

be broken and remade (alternative reaction pathway) by forming one or more intermediate compounds

The activation enthalpy needed to form intermediates from reactants, and then products from intermediates is lower than to form products directly from reactants

Reversible Reaction

Dynamic equilibrium – closed system – forward reaction going at same rate as backward reaction, amounts of products and reactants doesn’t change

Le Chatelier’s Principle: if there’s a change in concentration, pressure or temperature, the equilibrium will move to help counteract the change

Catalysts don’t affect position of equilibrium (yield) but do mean it is reached faster

Concentration Pressure TemperatureIncrease to right if more

reactant added, to left if more product added

shifts to side with fewest gas molecules to reduce pressure

moves in endothermic (+ΔH) to absorb extra heat

Decrease to left if reactant lowered, to right if product lowered

to side with most gas molecules to raise pressure

moves in exothermic (-ΔH) to replace heat

The Atmosphere The Earth’s surface emits much lower frequency radiation than Sun

(because Earth much cooler) Only molecules of different atoms absorb IR radiation because polarities

of their bonds change as they vibrate Carbon dioxide, water, methane etc do and they are greenhouse gases Gas molecules’ bonds have certain fixed vibrational energy levels so only

certain frequencies of radiation corresponding to these are absorbed UV and visible light move electrons up to higher quantised energy levels, if

enough energy absorbed the bonds break forming free radicals

The Greenhouse Effect IR window – range of IR frequencies that are not absorbed by atmospheric

gases Contribution of any particular gas to greenhouse effect depends on how

much radiation one molecule of gas absorbs, and how much of that gas there is in atmosphere

Climate change not new, natural – regular changes in Earth’s orbit around Dun cause ice age and interglacials; sunspot cycles; volcanic eruptions and meteor impacts cause cooling

Anthropogenic change is new and caused by humans Mass spectrometry shows composition of air inside ice in polar regions,

sea water is becoming more acidic (CO2 dissolves to form carbonic acid H2CO3)

Alternative fuels, fuel-efficient technologies, carbon capture and storage (injecting as liquid into deep ocean; storing under pressure deep underground; reacting with metal oxides to form carbonate minerals), increasing photosynthesis

Halogenoalkanes and CFCs

Bond fission – breaking a covalent bondo homolytic fission – one electron goes to each atom forming 2

electrically uncharged radicals (reactive particles due to unpaired electron)

o heterolytic fission – both electrons go to same atom, forming cation and anion

Halogens react with alkanes in photochemical reactions, forming Halogenoalkanes (free-radical substitution reaction)

o initiation reaction photodissociation of halogen molecule (homolytic fission)

o propagation reaction a radical attacks the alkane molecule the new alkane free radical can attack another halogen

molecule this happens until all halogen and alkane molecules used up

o termination reaction free radicals react to form stable molecules

CFCs are unreactive, non-flammable, harmless, range of boiling points so used for fire extinguishers, propellants in aerosols, coolant gas in fridges, expanding polystyrene for packaging and insulating material

HCFCs (still damage ozone layer but smaller effect) and HFCs – temporary alternatives until safer products developed (both are greenhouse gases, much worse than CO2)

Fridges – hydrocarbons, ammonia as coolant gas Aerosols – pump spray systems, nitrogen CO2 – foamed polymers

Ozone took time to make link between ozone depletion and Halogenoalkanes

The Polymer Revolution

Addition Reactions of Alkanes Hydrogen + ethane = ethane (nickel catalyst + 150 C + high pressure; or

platinum catalyst + rtp) Bromine water used to test for C=C bonds – shake alkene with orange

bromine water, the solution goes colourless (electrophilic addition reaction). If saturated compound, doesn’t react and solution stays brown

Electrophiles are electron-pair acceptors (positive ions and polar molecules) – attracted to double bonds

Alcohols and Other Organic Compounds Aldehydes have carbonyl group at end of carbon chain Ketones – anywhere except at end Carboxylic acid – COOH (carboxylic functional group (carbonyl and

hydroxyl)

Primary alcohols oxidise to aldehydes then carboxylic acids Secondary alcohols oxidise to ketones only Tertiary alcohols do not oxidise Oxidising agent acidified potassium dichromate(VI) (orange) is reduced to

green chromium(III) ion Cr3+ Oxidising primary alcohols

o gently heating ethanol with (VI) and sulphuric acid in test tube produces apple smelling ethanal

o to just get aldehyde, get it out of oxidising solution as soon as it’s formed by gently heating excess alcohol with controlled amount of oxidising agent in distillation apparatus, so aldehyde is distilled off immediately

o to produce carboxylic acid, alcohol has to be vigorously heated – mixed with excess oxidising agent and heated under reflux

Hydrogen Bonding Strongest intermolecular force O, F or N which are very electronegative draw bonding electrons away

from hydrogen, polarising bond, and H has high charge density In ice there is maximum number of hydrogen bonds, so less dense than

water, lots of big spaces in lattice structure Polymers dissolve in water if they can form hydrogen bonds with water

molecules instead of with each other If too many OH groups hydrogen bonding is too strong, meaning too much

energy needed to break it down – insoluble If too few OH groups, won’t be many hydrogen bonds formed with water –

insoluble

Polymers Copolymers made from more than one type of monomer joined in a

random order Thermoplastic polymers – no cross-linking, chains held together by weak

intermolecular forces which are easy to overcome, easy to melt, hardens into new shape when it cools, can be remoulded

Thermosetting polymers – covalent cross-links which hold chains together in 3D giant covalent structure, doesn’t soften when heated but chars, strong, hard, rigid and insoluble

LDPE – soft, flexible, used for plastic bags, squeezy bottles Poly(propene) – tough, strong, used for bottle crates, rope Poly(chloroethene) (PVC) – durable, flexible, used for water pipes,

insulation on electric wires, costumes, as a building material PTFE – inert and non-stick properties so used as frying pan coating Polystyrene – cheap and can be made into expanded polystyrene which is

light and good insulator – disposable cups, crash helmets Perspex – transparent and strong so can replace glass

E/Z Isomerism Atoms can’t rotate around double bonds Stereoisomers have the same structural formula but a different

arrangement in space Trans/E-isomer is when different groups are across double bond

(opposite) Cis/Z-isomer is when same groups are both above or below double bond

(together) Br higher priority than F CH3 higher priority than H

Infrared Spectroscopy A beam of IR radiation is passed through sample of chemical which is

absorbed by covalent bonds, increasing their vibrational energy Bonds between different atoms and bonds in different places in a

molecule absorb different frequencies of IR radiation Fingerprint region (1000 to 1550 cm-1) identifies a molecule because it is

unique to a particular compound